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seleniumBrominekrypton
Cl

Br

I
Appearance
gas/liquid: red-brown
solid: metallic luster
General properties
Name, symbol, number Bromine, Br, 35
Element category halogen
Group, period, block 174, p
Standard atomic weight 79.904(1)g·mol−1
Electron configuration [Ar] 4s2 3d10 4p5
Electrons per shell 2, 8, 18, 7 (Image)
Physical properties
Phase liquid
Density (near r.t.) (Br2, liquid) 3.1028 g·cm−3
Melting point 265.8 K, -7.2 °C, 19 °F
Boiling point 332.0 K, 58.8 °C, 137.8 °F
Critical point 588 K, 10.34 MPa
Heat of fusion (Br2) 10.571 kJ·mol−1
Heat of vaporization (Br2) 29.96 kJ·mol−1
Specific heat capacity (25 °C) (Br2)
75.69 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 185 201 220 244 276 332
Atomic properties
Oxidation states 7, 5, 4, 3, 1, -1
(strongly acidic oxide)
Electronegativity 2.96 (Pauling scale)
Ionization energies 1st: 1139.9 kJ·mol−1
2nd: 2103 kJ·mol−1
3rd: 3470 kJ·mol−1
Atomic radius 120 pm
Covalent radius 120±3 pm
Van der Waals radius 185 pm
Miscellanea
Crystal structure orthorhombic
Magnetic ordering diamagnetic[1]
Electrical resistivity (20 °C) 7.8×1010Ω·m
Thermal conductivity (300 K) 0.122 W·m−1·K−1
Speed of sound (20°C) 206 m/s
CAS registry number 7726-95-6
Most stable isotopes
Main article: Isotopes of Bromine
iso NA half-life DM DE (MeV) DP
79Br 50.69% 79Br is stable with 44 neutrons
81Br 49.31% 81Br is stable with 46 neutrons

Bromine (pronounced /ˈbroʊmiːn/ BROH-meen or /ˈbroʊmɨn/ BROH-min, from Greek: βρῶμος, brómos, meaning "stench (of he-goats)"),[2] is a chemical element with the symbol Br and atomic number 35. A halogen element, bromine is a reddish-brown volatile liquid at standard room temperature that is intermediate in reactivity between chlorine and iodine. Bromine vapors are corrosive and toxic. Approximately 556,000 metric tons were produced in 2007.[3] The main applications for bromine are in fire retardants and fine chemicals.

Contents

History

Illustrative and secure bromine sample for teaching

Bromine was discovered independently by two chemists Antoine Balard[4] and Carl Jacob Löwig[5] in 1825 and 1826.[6]

Balard found bromide chemicals in the ash of sea weed from the salt marshes of Montpellier in 1826. The sea weed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance resembled that of an intermediate of chlorine and iodine; with those results he tried to prove that the substance was iodine monochloride (ICl), but after failing to do so he was sure that he had found a new element and named it muride, derived from the Latin word muria for brine.[4]

Carl Jacob Löwig isolated bromine from a mineral water spring from his hometown Bad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with diethylether. After evaporation of the ether a brown liquid remained. With this liquid as a sample for his work he applied for a position in the laboratory of Leopold Gmelin in Heidelberg. The publication of the results was delayed and Balard published his results first.[5]

After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique.[7] In his publication Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors.[8] Bromine was not produced in large quantities until 1860.

The first commercial use, besides some minor medical applications, was the use of bromine for the daguerreotype. In 1840 it was discovered that bromine had some advantages over the previously used iodine vapor to create the light sensitive silver halide layer used for daguerreotypy.[9]

Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, until they were gradually superseded by chloral hydrate and then the barbiturates.[10]

Characteristics

Bromine is the only nonmetallic element that is liquid at room temperature, and one of only two elements on the periodic table that are liquid at room temperature (mercury is the other). The melting point of bromine is −7.2 °C and the boiling point 58.8 °C (138 °F). The pure chemical element has the physical form of a diatomic molecule, Br2. It is a dense, mobile, slightly transparent reddish-brown liquid, that evaporates easily at standard temperature and pressures to give a red vapor (its color resembles nitrogen dioxide) that has a strong disagreeable odor resembling that of chlorine. Bromine is a halogen, and is less reactive than chlorine and more reactive than iodine. Bromine is slightly soluble in water, and highly soluble in carbon disulfide, aliphatic alcohols (such as methanol), and acetic acid. It bonds easily with many elements and has a strong bleaching action. Bromine, like chlorine, is also used in maintenance of swimming pools.

Certain bromine-related compounds have been evaluated to have an ozone depletion potential or bioaccumulate in living organisms. As a result many industrial bromine compounds are no longer manufactured, are being restricted, or scheduled for phasing out. The Montreal Protocol mentions several organobromine compounds for this phase out.

Bromine is a powerful oxidizing agent. It reacts vigorously with metals, especially in the presence of water, as well as most organic compounds, especially upon illumination.

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Isotopes

Bromine has 2 stable isotopes: 79Br (50.69 %) and 81Br (49.31%). At least another 23 radioisotopes are known to exist.[11] Many of the bromine isotopes are fission products. Several of the heavier bromine isotopes from fission are delayed neutron emitters. All of the radioactive bromine isotopes are relatively short lived. The longest half life is the neutron deficient 77Br at 2.376 days. The longest half life on the neutron rich side is 82Br at 1.471 days. A number of the bromine isotopes exhibit metastable isomers. Stable 79Br exhibits a radioactive isomer, with a half life of 4.86 seconds. It decays by isomeric transition to the stable ground state.[12]

Allotropes

At a pressure of 55 GPa bromine converts to a metal. At 75 GPa it converts to a face centered orthorhombic structure. At 100 GPa it converts to a body centered orthorhombic monoatomic form.[13]

Occurrence and production

World bromine production trend
View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine 31°9′0″N 35°27′0″E / 31.15°N 35.45°E / 31.15; 35.45

The diatomic element Br2 does not occur naturally. Instead, bromine exists exclusively as bromide salts in diffuse amounts in crustal rock. Due to leaching, bromide salts have accumulated in sea water (65 ppm),[14] but at a lower concentration than chloride. Bromine may be economically recovered from bromide-rich brine wells and from the Dead Sea waters (up to 50000 ppm).[15][16]

Approximately 556,000 metric tonnes (worth around US$2.5 billion) of bromine are produced per year (2007) worldwide with the United States, Israel, and China being the primary producers.[17][18][19] Bromine production has increased sixfold since the 1960s. The largest bromine reserve in the United States is located in Columbia and Union County, Arkansas, U.S.[20] China's bromine reserves are located in the Shandong Province and Israel's bromine reserves are contained in the waters of the Dead Sea. The bromide-rich brines are treated with chlorine gas, flushing through with air. In this treatment, bromide anions are oxidized to bromine by the chlorine gas.

2 Br + Cl2 → 2 Cl + Br2

Because of its commercial availability and long shelf-life, bromine is not typically prepared. Small amounts of bromine can however be generated through the reaction of solid sodium bromide with concentrated sulfuric acid (H2SO4). The first stage is formation of hydrogen bromide (HBr), which is a gas, but under the reaction conditions some of the HBr is oxidized further by the sulfuric acid to form bromine (Br2) and sulfur dioxide (SO2).

NaBr (s) + H2SO4 (aq) → HBr (aq) + NaHSO4 (aq)
2 HBr (aq) + H2SO4 (aq) → Br2 (g) + SO2 (g) + 2 H2O (l)

Similar alternatives, such as the use of dilute hydrochloric acid with sodium hypochlorite, are also available. The most important thing is that the anion of the acid (in the above examples, sulfate and chloride, respectively) be more electronegative than bromine, allowing the substitution reaction to occur.

Reaction involving a strong oxidizing agent, such as potassium permanganate, on bromide ions in the presence of an acid also gives bromine. An acidic solution of bromate ions and bromide ions will also disproportionate slowly to give bromine.

Bromine is only slightly soluble in water. But the solubility can be increased by the presence of bromide ions. However, concentrated solutions of bromine are rarely prepared in the lab as they will continually give off toxic red-brown bromine gas due to its very high vapor pressure. Sodium thiosulphate is an excellent reagent for getting rid of bromine completely including the stains and odor.

Compounds

Organic chemistry

N-Bromosuccinimide

Organic compounds are brominated by either addition or substitution reactions. Bromine undergoes electrophilic addition to the double-bonds of alkenes, via a cyclic bromonium intermediate. In non-aqueous solvents such as carbon disulfide, this affords the di-bromo product. For example, reaction with ethylene will produce 1,2-dibromoethane. Bromine also undergoes electrophilic addition to phenols and anilines. When used as bromine water, a small amount of the corresponding bromohydrin is formed as well as the dibromo compound. So reliable is the reactivity of bromine that bromine water is employed as a reagent to test for the presence of alkenes, phenols, and anilines. Like the other halogens, bromine participates in free radical reactions. For example hydrocarbons are brominated upon treatment with bromine in the presence of light.

Bromine, sometimes with a catalytic amount of phosphorus, easily brominates carboxylic acids at the α-position. This method, the Hell-Volhard-Zelinsky reaction, is the basis of the commercial route to bromoacetic acid. N-Bromosuccinimide is commonly used as a substitute for elemental bromine, being easier to handle, and reacting more mildly and thus more selectively. Organic bromides are often preferable relative to the less reactive chlorides and more expensive iodide-containing reagents. Thus, Grignard and organolithium compound are most often generated from the corresponding bromides.

Inorganic chemistry

Oxidation states
of bromine
-1 HBr
0 Br2
+1 BrCl
+3 BrF3
+5 BrF5
+5 BrO3
+7 BrO4

Bromine is an oxidizer, and it will oxidize iodide ions to iodine, being itself reduced to bromide:

Br2 + 2 I → 2 Br + I2

Bromine will also oxidize metals and metalloids to the corresponding bromides. Anhydrous bromine is less reactive toward many metals than hydrated bromine, however. Dry bromine reacts vigorously with aluminium, titanium, mercury as well as alkaline earths and alkali metals.

If bromine is dissolved in hydroxide containing water not only bromide (Br) is formed, but also the hypobromite (OBr). This hypobromite is responsible for the bleaching abilities of bromide solutions. In warm solutions the disproportion reaction of the hypobromite is quantitative. The resulting bromate is a strong oxidising agent and very similar to the chlorate.

3 BrOBrO3 + 2 Br

The perbromates are not accessible through electrolysis like the perchlorates, but only by reacting bromate solutions with fluorine or ozone.

BrO3 + H2O + F2BrO4 + 2 HF
BrO3 + O3BrO4 + O2

Applications

A wide variety of organobromine compounds are used in industry. Some are prepared from bromine and others are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine.[3]

Illustrative of the addition reaction[21] is the preparation of 1,2-Dibromoethane, the organobromine compound produced in the largest amounts:

C2H4 + Br2 → CH2BrCH2Br

Flame retardant

Tetrabromobisphenol A

Brominated flame retardants represent a commodity of growing importance. If the material burns the flame retardants produce hydrobromic acid which interferes in the radical chain reaction of the oxidation reaction of the fire. The highly reactive hydrogen oxygen and hydroxy radicals react with hydrobromic acid and form less reactive bromine radicals.[22][23] The bromine containing compounds can be placed in the polymers either during polymerization if a small amount of brominated monomer is added or the bromine containing compound is added after polymerization. Tetrabromobisphenol A can be added to produce polyesters or epoxy resins. Epoxy used in printed circuit boards (PCB) are normally made from flame retardant resins, indicated by the FR in the abbreviation of the products (FR-4 and FR-2. Vinyl bromide can be used in the production of polyethylene, polyvinylchloride or polypropylene. Decabromodiphenyl ether can be added to the final polymers.[24]

Gasoline additive

Ethylene bromide was an additive in gasolines containing lead anti-engine knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% of the bromine uses in 1966 in the US. This application has declined since the 1970s due to environmental regulations.[25] Ethylene bromide is also used as a fumigant, but again this application is declining.[19]

Pesticide

Methyl bromide (bromomethane)

Methyl bromide was widely used as pesticide to fumigate soil. The Montreal Protocol on Substances that Deplete the Ozone scheduled the phase out for the ozone depleting chemical until 2005. In 1991, an estimated 35,000 metric tonnes of the chemical were used to control nematodes, fungi, weeds and other soil-borne diseases.[26][27]

Medical and veterinary

  • Bromide compounds, especially potassium bromide, were frequently used as sedatives in the 19th and early 20th century. Bromides in the form of simple salts are still used as anticonvulsants in both veterinary and human medicine.

Other uses

Orange fluoresces of DNA Ethidium bromide intercalate
  • The bromides of calcium, sodium, and zinc account for a sizable part of the bromine market. These salts form dense solutions in water that are used as drilling fluids sometimes called clear brine fluids.[19][28]
  • Bromine is also used in the production of brominated vegetable oil, which is used as an emulsifier in many citrus-flavored soft drinks (e.g. Mountain Dew). After the introduction in the 1940s the compound was extensively used until the UK and the US limited its use in the mid 1970s and alternative emulsifiers were developed.[29]
    Soft drinks containing brominated vegetable oil are still sold in the US (2009).[30]
Tralomethrin
  • Several dyes, agrichemicals, and pharmaceuticals are organobromine compounds. 1-Bromo-3-chloropropane, 1-bromoethylbenzene, and 1-bromoalkanes are prepared by the antimarkovnikov addition of HBr to alkenes. Ethidium bromide, EtBr, is used as a DNA stain in gel electrophoresis.
  • High refractive index compounds
  • Water purification compounds, disinfectants and insecticides, such as tralomethrin (C22H19Br4NO3).[19]
  • Potassium bromide is used in some photographic developers to inhibit the formation of fog (undesired reduction of silver).
  • Vapor is used as the second step in sensitizing daguerreotype plates to be developed under Mercury (Hg) vapor. Bromine acts as an accelerator to the light sensitivity of the previously iodized plate.
  • Bromine is also used to reduce mercury pollution from coal-fired power plants. This can be achieved either by treating activated carbon with bromine or by injecting bromine compounds onto the coal prior to combustion.

Biological role

Bromine has no known essential role in human or mammalian health, but inorganic bromine and organobromine compounds do occur naturally, and some may be of use to higher organisms in dealing with parasites. For example, in the presence of H2O2 formed by the eosinophil, and either chloride or bromide ions, eosinophil peroxidase provides a potent mechanism by which eosinophils kill multicellular parasites (such as, for example, the nematode worms involved in filariasis); and also certain bacteria (such as tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that preferentially uses bromide over chloride for this purpose, generating hypobromite (hypobromous acid).[31]

Marine organisms are the main source of organobromine compounds. Over 1600 compounds were identified by 1999. The most abundant one is methyl bromide with estimated 56,000 metric tonnes produced by marine algae.[32] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% methyl bromide.[33] A famous example of a bromine-containing organic compound that has been used by humans for a long time is Tyrian purple.[32][34] The brominated indigo is produced by a medium-sized predatory sea snail, the marine gastropod Murex brandaris. It took until 1909 before the organobromine nature of the compound was discovered by Paul Friedländer.[35] Most organobromine compounds in nature arise via the action of vanadium bromoperoxidase.[36]

Safety

Elemental bromine is toxic and causes burns. As an oxidizing agent, it is incompatible with most organic and inorganic compounds. Care needs to be taken when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames.

When certain ionic compounds containing bromine are mixed with potassium permanganate (KMnO4) and an acidic substance, they will form a pale brown cloud of bromine gas. This gas smells like bleach and is very irritating to the mucus membranes. Upon exposure, one should move to fresh air immediately. If symptoms of bromine poisoning arise, medical attention is needed.

References

  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81st edition, CRC press.
  2. ^ Gemoll W, Vretska K (1997). Griechisch-Deutsches Schul- und Handwörterbuch ("Greek-German dictionary"), 9th ed.. öbvhpt. ISBN 3-209-00108-1. 
  3. ^ a b Jack F. Mills (2002). Bromine: in Ullmann's Encyclopedia of Chemical Technology. Weinheim: Wiley-VCH Verlag. doi:10.1002/14356007.a04_391. 
  4. ^ a b Balard, Antoine (1826). "Memoire of a peculire Substance contained in Sea Water". Annals of Philosophy: 387– and 411–. http://books.google.com/books?id=A-M4AAAAMAAJ. 
  5. ^ a b Landolt, Hans Heinrich (1890). "Nekrolog: Carl Löwig". Berichte der deutschen chemischen Gesellschaft 23: 905. doi:10.1002/cber.18900230395. http://gallica.bnf.fr/ark:/12148/bpt6k907222/f920.chemindefer. 
  6. ^ Weeks, Mary Elvira (1932). "The discovery of the elements: XVII. The halogen family.". Journal of Chemical Education 9: 1915. 
  7. ^ Balard, A.J. (1826). Annales de Chimie et Physique 32: 337. 
  8. ^ Wisniak, Jaime (2004). "Antoine-Jerôme Balard. The discoverer of bromine". Revista CENIC Ciencias Químicas 35. http://revistas.mes.edu.cu:9900/EDUNIV/03-Revistas-Cientificas/Rev.CENIC-Ciencias-Quimicas/2004/1/09204109.pdf. 
  9. ^ Barger, M. Susan; White, William Blaine (2000). "Technological Practice of Daguerreotypy". The Daguerreotype: Nineteenth-century Technology and Modern Science. JHU Press. pp. 31–35. ISBN 9780801864582. 
  10. ^ Shorter, Edward (1997). A History of Psychiatry: From the Era of the Asylum to the Age of Prozac. John Wiley and Sons. p. 200. ISBN 9780471245315. 
  11. ^ GE (1989). Chart of the Nuclides, 14th Edition. Nuclear Energy. 
  12. ^ Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3. doi:10.1016/j.nuclphysa.2003.11.001. 
  13. ^ Duan, Defang (2007-09-26). "Ab initio studies of solid bromine under high pressure". Pysical Review B 76. doi:10.1103/PhysRevB.76.104113. 
  14. ^ Tallmadge, John A.; Butt, John B.; Solomon Herman J. (1964). "Minerals From Sea Salt". Ind. Eng. Chem. 56: 44. doi:10.1021/ie50655a008. 
  15. ^ Oumeish, Oumeish Youssef (1996). "Climatotherapy at the Dead Sea in Jordan". Clinics in Dermatology 14: 659. doi:10.1016/S0738-081X(96)00101-0. 
  16. ^ Al-Weshah, Radwan A. (2008). "The water balance of the Dead Sea: an integrated approach". Hydrological Processes 14: 145. doi:10.1002/(SICI)1099-1085(200001)14:1<145::AID-HYP916>3.0.CO;2-N. 
  17. ^ Emsley, John (2001). "Bromine". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 69–73. ISBN 0198503407. 
  18. ^ Lyday, Phyllis A.. "Comodity Report 2007: Bromine". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/bromine/bromimcs07.pdf. Retrieved 2008-09-03. 
  19. ^ a b c d Lyday, Phyllis A.. "Mineral Yearbook 2007: Bromine". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/bromine/myb1-2006-bromi.pdf. Retrieved 2008-09-03. 
  20. ^ "Bromine:An Important Arkansas Industry". Butler Center for Arkansas Studies. http://www.cals.lib.ar.us/butlercenter/lesson_plans/lesson%20plans/Lesson%20plans-retained/Bromine.pdf. 
  21. ^ N. A. Khan, F. E. Deatherage, and J. B. Brown (1963), "Stearolic Acid", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV4P0851 ; Coll. Vol. 4: 851 
  22. ^ Green, Joseph (1996). "Mechanisms for Flame Retardancy and Smoke suppression -A Review". Journal of Fire Sciences 14: 426. doi:10.1177/073490419601400602. 
  23. ^ Kaspersma, Jelle; Doumena, Cindy; Munrob Sheilaand; Prinsa, Anne-Marie (2002). "Fire retardant mechanism of aliphatic bromine compounds in polystyrene and polypropylene". Polymer Degradation and Stability 77: 325. doi:10.1016/S0141-3910(02)00067-8. 
  24. ^ Weil, Edward D.; Levchik, Sergei (2004). "A Review of Current Flame Retardant Systems for Epoxy Resins". Journal of Fire Sciences 22: 25. doi:10.1177/0734904104038107. 
  25. ^ Alaeea, Mehran; Ariasb, Pedro; Sjödinc, Andreas; Bergman, Åke (2003). "An overview of commercially used brominated flame retardants, their applications, their use patterns in different countries/regions and possible modes of release". Environment International 29: 683. doi:10.1016/S0160-4120(03)00121-1. 
  26. ^ Messenger, Belinda; Braun, Adolf (2000). "Alternatives to Methyl Bromide for the Control of Soil-Borne Diseases and Pests in California". Pest Management Analysis and Planning Program. http://www.cdpr.ca.gov/docs/emon/methbrom/alt-anal/sept2000.pdf. Retrieved 2008-11-17. 
  27. ^ Decanio, Stephen J.; Norman, Catherine S. (2008). "Economics of the "Critical Use" of Methyl bromide under the Montreal Protocol". Contemporary Economic Policy 23: 376. doi:10.1093/cep/byi028. 
  28. ^ Darley, H. C. H.; Gray, George Robert (1988). Composition and Properties of Drilling and Completion Fluids. Gulf Professional Publishing. pp. 61–62. ISBN 9780872011472. 
  29. ^ Kaufman, Vered R.; Garti, Nissim (1984). "Effect of cloudy agents on the stability and opacity of cloudy emulsions for soft drinks". International Journal of Food Science & Technology 19: 255. doi:10.1111/j.1365-2621.1984.tb00348.x. 
  30. ^ Horowitz, B. Zane (1997). "Bromism from Excessive Cola Consumption',Clinical Toxicology". Clinical Toxicology 35: 315. doi:10.3109/15563659709001219. 
  31. ^ Mayeno AN, Curran AJ, Roberts RL, Foote CS (April 1989). "Eosinophils preferentially use bromide to generate halogenating agents". J. Biol. Chem. 264 (10): 5660–8. PMID 2538427. 
  32. ^ a b Gordon W. Gribble (1999). "The diversity of naturally occurring organobromine compounds". Chemical Society Reviews 28: 335. doi:10.1039/a900201d. 
  33. ^ Burreson, B. Jay; Moore, Richard E.; Roller, Peter P. (1976). "Volatile halogen compounds in the alga Asparagopsis taxiformis (Rhodophyta)". Journal of Agricultural snd Food Chemistry 24: 856. doi:10.1021/jf60206a040. 
  34. ^ Gordon W. Gribble (1998). "Naturally Occurring Organohalogen Compounds". Acc. Chem. Res. 31: 141. doi:10.1021/ar9701777. 
  35. ^ Friedländer, P. (1909). "Über den Farbstoff des antiken Purpurs aus murex brandaris". Berichte der deutschen chemischen Gesellschaft 42: 765. doi:10.1002/cber.190904201122. 
  36. ^ Butler, Alison; Carter-Franklin, Jayme N. (2004). "The role of vanadium bromoperoxidase in the biosynthesis of halogenated marine natural products". Natural Product Reports 21: 180. doi:10.1039/b302337k. 

External links


1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

BROMINE (symbol Br, atomic weight 79-96), a chemical element of the halogen group, which takes its name from its pungent unpleasant smell (0pW,uos, a stench). It was first isolated by A. J. Balard in 1826 from the salts in the waters of the Mediterranean. He established its elementary character, and his researches were amplified by K. Lbwig (1803-1890) in Das Brom and seine chemischen Verhaltnisse (1829). Bromine does not occur in nature in the uncombined condition, but in combination with various metals is very widely but sparingly distributed. Potassium, sodium and magnesium bromides are found in mineral waters, in river and sea-water, and occasionally in marine plants and animals. Its chief commercial sources are the salt deposits at Stassfurt in Prussian Saxony, in which magnesium bromide is found associated with various chlorides, and the brines of Michigan, Ohio, Pennsylvania and West Virginia, U.S.A.; small quantities are obtained from the mother liquors of Chile saltpetre and kelp. In combination with silver it is found as the mineral bromargyrite (bromite).

Table of contents

Manufacture

The chief centres of the bromine industry are Stassfurt and the central district of Michigan. It is manufactured from the magnesium bromide contained in "bittern" (the mother liquor of the salt industry), by two processes, the continuous and the periodic. The continuous process depends upon the decomposition of the bromide by chlorine, which is generated in special stills. A regular current of chlorine mixed with steam is led in at the bottom of a tall tower filled with broken bricks, and there meets a descending stream of hot bittern: bromine is liberated and is swept out of the tower together with some chlorine, by the current of steam, and then condensed in a worm. Any uncondensed bromine vapour is absorbed by moist iron borings, and the resulting iron bromide is used for the manufacture of potassium bromide. The periodic process depends on the interaction between manganese dioxide (pyrolusite), sulphuric acid, and a bromide, and the operation is carried out in sandstone stills heated to 60° C., the product being condensed as in the continuous process. The substitution of potassium chlorate for pyrolusite is recommended when calcium chloride is present in the bittern. The crude bromine is purified by repeated shaking with potassium, sodium or ferrous bromide and subsequent redistillation. Commercial bromine is rarely pure, the chief impurities present in it being chlorine, hydrobromic acid, and bromoform (M. Hermann, Annalen, 18 55, 95, p. 211). E. Gessner (Berichte, 1876, 9, p. 1507) removes chlorine by repeated shaking with water, followed by distillation over sulphuric acid; hydrobromic acid is removed by distillation with pure manganese dioxide, or mercuric oxide, and the product dried over sulphuric acid. J. S. Stas, in his stoichiometric researches, prepared chemically pure bromine from potassium bromide, by converting it into the bromate which was purified by repeated crystallization. By heating the bromate it was partially converted into the bromide, and the resulting mixture was distilled with sulphuric acid. The distillate was further purified by digestion with milk of lime, precipitation with water, and further digestion with calcium bromide and barium oxide, and was finally redistilled.

Characters

Bromine at ordinary temperatures is a mobile liquid of fine red colour, which appears almost black in thick layers. It boils at 59° C. According to Sir W. Ramsay and S. Young, bromine, when dried over sulphuric acid, boils at 57.65° C., and when dried over phosphorus pentoxide, boils at 58.85° C. (under a pressure of 755.8 mm.), forming a deep red vapour, which exerts an irritating and directly poisonous action on the respiratory organs. It solidifies at - 21° C. (Quincke) to a dark brown solid. Its specific gravity is 3.18828 (r), latent heat of fusion 16.185 calories, latent heat of vaporization 45.6 calories, specific heat 0.1071. The specific heat of bromine vapour, at constant pressure, is 0.05504 and at constant volume is 0.04251 (K. Strecker). Bromine is soluble in water, to the extent of 3.226 grammes of bromine per too grammes of solution at 15° C., the solubility being slightly increased by the presence of potassium bromide. The solution is of an orange-red colour, and is quite permanent in the dark, but on exposure to light, gradually becomes colourless, owing to decomposition into hydrobromic acid and oxygen. By cooling the aqueous solution, hyacinth-red octahedra of a crystalline hydrate of composition Br 4H 2 O or Br2.8H20 are obtained (Bakhuis Roozeboom, Zeits. phys. Chem., 1888, 2. P. 449). Bromine is readily soluble in chloroform, alcohol and ether.

Its chemical properties are in general intermediate between those of chlorine and iodine; thus it requires the presence of a catalytic agent, or a fairly high temperature, to bring about its union with hydrogen. It does not combine directly with oxygen, nitrogen or carbon. With the other elements it unites to form bromides, often with explosive violence; phosphorus detonates in liquid bromine and inflames in the vapour; iron is occasionally used to absorb bromine vapour, potassium reacts energetically, but sodium requires to be heated to 200° C. The chief use of bromine in analytical chemistry is based upon the oxidizing action of bromine water. Bromine and bromine water both bleach organic colouring matters.

FIG.. - Fruit of the pine-apple (Ananas saliva), consisting of numerous flowers and bracts united together so as to form a collective or anthocarpous fruit. The crown of the pine-apple, c, consists of a series of empty bracts prolonged beyond the fruit.

The use of bromine in the extraction of gold was proposed by R. Wagner (Dingier's Journal, 218, p. 253) and others, but its cost has restricted its general application. Bromine is used extensively in organic chemistry as a substituting and oxidizing agent and also for the preparation of addition compounds. Reactions in which it is used in the liquid form, in vapour, in solution, and in the presence of the so-called "bromine carriers." have been studied. Sunlight affects the action of bromine vapour on organic compounds in various ways, sometimes retarding or accelerating the reaction, while in some cases the products are different (J. Schramm, Monatshefte fiir Chemie, 1887, 8, p. tot). Some reactions, which are only possible by the aid of nascent bromine, are carried out by using solutions of sodium bromide and bromate, with the amount of sulphuric acid calculated according to the equation 5NaBr NaBr03-1-6H2S04= 6NaHSO 4 3H 2 O 6Br. (German Patent, 26642.) The diluents in which bromine is employed are usually ether, chloroform, acetic acid, hydrochloric acid, carbon bisulphide and water, and, less commonly, alcohol, potassium bromide and hydrobromic acid; the excess of bromine being removed by heating, by sulphurous acid or by shaking with mercury. The choice of solvent is important, for the velocity of the reaction and the nature of the product may vary according to the solvent used, thus A. Baeyer and F. Blom found that on brominating orthoacetamido-acetophenone in presence of water or acetic acid, the bromine goes into the benzene nucleus, whilst in chloroform or sulphuric acid or by use of bromine vapour it goes into the side chain as well. The action of bromine is sometimes accelerated by the use of compounds which behave catalytically, the more important of these substances being iodine, iron, ferric chloride, ferric bromide, aluminium bromide and phosphorus. For oxidizing purposes bromine is generally employed in aqueous and in alkaline solutions, one of its most important applications being by Emil Fischer (Berichte, 1889, 22, p. 362) in his researches on the sugars. The atomic weight of bromine has been determined by J. S. Stas and C. Marignac from the analysis of potassium bromide, and of silver bromide. G. P. Baxter (Zeit. anorg. Chem. 1906, 50, p. 389) determined the ratios Ag: AgBr, and AgC1: Ag Br.

Hydrobromic Acid

This acid, HBr, the only compound of hydrogen and bromine, is in many respects similar to hydrochloric acid, but is rather less stable. It may be prepared by passing hydrogen gas and bromine vapour through a tube containing a heated platinum spiral. It cannot be prepared with any degree of purity by the action of concentrated sulphuric acid on bromides, since secondary reactions take place, leading to the liberation of free bromine and formation of sulphur dioxide. The usual method employed for the preparation of the gas consists in dropping bromine on to a mixture of amorphous phosphorus and water, when a violent reaction takes place and the gas is rapidly liberated. It can be obtained also, although in a somewhat impure condition, by the direct action of bromine on various saturated hydrocarbons (e.g. paraffin-wax), while an aqueous solution may be obtained by passing sulphuretted hydrogen through bromine water. Alexander Scott (Journal of Chem. Soc., 1900, 77, p. 648) prepares pure hydrobromic acid by covering bromine, which is contained in a large flask, with a layer of water, and passing sulphur dioxide into the water above the surface of the bromine, until the whole is of a pale yellow colour; the resulting solution is then distilled in a slow current of air and finally purified by distillation over barium bromide. At ordinary temperatures hydrobromic acid is a colourless gas which fumes strongly in moist air, and has an acid taste and reaction. It can be condensed to a liquid, which boils at - 64.9°C. (under a pressure of 738.2 mm.), and, by still further cooling, gives colourless crystals which melt at - 88.5° C. It is readily soluble in water, forming the aqueous acid, which when saturated at 0° C. has a specific gravity of 1.78. When boiled, the aqueous acid loses either acid or water until a solution of constant boiling point is obtained, containing 48% of the acid and boiling at 126° C. under atmospheric pressure; should the pressure, however, vary, the strength of the solution boiling at a constant temperature varies also. Hydrobromic acid is one of the "strong" acids, being ionized to a very large extent even in concentrated solution, as shown by the molecular conductivity increasing by only a small amount over a wide range offdilution.

Bromides

Hydrobromic acid reacts with metallic oxides, hydroxides and carbonates to form bromides, which can in many cases be obtained also by the direct union of the metals with bromine. As a class, the metallic bromides are solids at ordinary temperatures, which fuse readily and volatilize on heating. The majority are soluble in water, the chief exceptions being silver bromide, mercurous bromide, palladious bromide and lead bromide; the last is, however, soluble in hot water. They are decomposed by chlorine, with liberation of bromine and formation of metallic chlorides; concentrated sulphuric acid also decomposes them, with formation of a metallic sulphate and liberation of bromine and sulphur dioxide. The non-metallic bromides are usually liquids, which are readily decomposed by water. Hydrobromic acid and its salts can be readily detected by the addition of chlorine water to their aqueous solutions, when bromine is liberated; or by warming with concentrated sulphuric acid and manganese dioxide, the same result being obtained. Silver nitrate in the presence of nitric acid gives with bromides a pale yellow precipitate of silver bromide, AgBr, which is sparingly soluble in ammonia. For their quantitative determination they are precipitated in nitric acid solution by means of silver nitrate, and the silver bromide well washed, dried and weighed.

No oxides of bromine have as yet been isolated, but three oxy-acids are known, namely hypobromous acid, HBrO, bromous acid, HBr02, and bromic acid, HBrO 3. Hypobromous acid is obtained by shaking together bromine water and precipitated mercuric oxide, followed by distillation of the dilute solution in vacuo at low temperature (about 40°C.). It is a very unstable compound, breaking up, on heating, into bromine and oxygen. The aqueous solution is light yellow in colour, and possesses strong bleaching properties. Bromous acid is formed by adding bromine to a saturated solution of silver nitrate (A. H. Richards, J. Soc. Chem. Ind., 1906, 25, p. 4). Bromic acid is obtained by the addition of the calculated amount of sulphuric acid (previously diluted with water) to the barium salt; by the action of bromine on the silver salt, in the presence of water, 5AgBrO, 3Br 2 3H 2 O = 5AgBr 6HBrO 3, or bypassing chlorine through asolution of bromine in water. The acid is only known in the form of its aqueous solution; this is, however, very unstable, decomposing on being heated to 100° C. into water, oxygen and bromine. By reducing agents such, for example, as sulphuretted hydrogen and sulphur-dioxide, it is rapidly converted into hydrobromic acid. Hydrobromic acid decomposes it according to the equation HBrO, 5HBr=3H20 3Br2. Its salts are known as bromates, and are as a general rule difficultly soluble in water, and decomposed by heat, with evolution of oxygen.

Applications

The salts of bromine are widely used in photography, especially bromide of silver. For antiseptic purposes it has been prepared as "bromum solidificatum," which consists of kieselguhr or similar substance impregnated with about 75% of its weight of bromine. In medicine it is largely employed in the form of bromides of potassium, sodium and ammonium, as well as in combination with alkaloids and other substances.

Medicinal Use. - Bromide of potassium is the safest and most generally applicable sedative of the nervous system. Whilst very weak, its action is perfectly balanced throughout all nervous tissue, so much so that Sir Thomas Lauder Brunton has suggested its action to be due to its replacement of sodium chloride (common salt) in the fluids of the nervous system. Hence bromide of potassium - or bromide of sodium, which is possibly somewhat safer still though not quite so certain in its action - is used as a hypnotic, as the standard anaphrodisiac, as a sedative in mania and all forms of morbid mental excitement, and in hyperaesthesia of all kinds. Its most striking success is in epilepsy, for which it is the specific remedy. It may be given in doses of from ten to fifty grains or more, and may be continued without ill effect for long periods in grave cases of epilepsy (grand mal). Of the three bromides in common use the potassium salt is the most rapid and certain in its action, but may depress the heart in morbid states of that organ; in such cases the sodium salt - of which the base is inert - may be employed. In whooping-cough, when a sedative is required but a stimulant is also indicated, ammonium bromide is often invaluable. The conditions in which bromides are most frequently used are insomnia, epilepsy, whooping-cough, delirium tremens, asthma, migraine, laryngismus stridulus, the symptoms often attendant upon the climacteric in women, hysteria, neuralgia, certain nervous disorders of the heart, strychnine poisoning, nymphomania and spermatorrhoea. Hydrobromic acid is often used to relieve or prevent the headache and singing in the ears that may follow the administration of quinine and of salicylic acid or salicylates.


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Simple English

Bromine (Br) is a chemical element. Its atomic number (which is the number of protons in it) is 35, and its atomic weight is 80. It is part of the Group 7 elements (halogens) on the periodic table. It is diatomic, which means that two atoms are stuck together to make a molecule in any bromine sample.

Contents

Properties

Physical properties

Bromine is a red-brown liquid. Liquid elements are rare; only bromine and mercury are liquid at room temperature. It easily evaporates to make suffocating brown fumes. It has a bad smell. Its name means "stench of he-goats". It can become a metal at very high pressures.

It has two stable isotopes. They are 79Br and 81Br. There are about 29 other radioactive isotopes.

Chemical properties

Bromine is quite reactive. Its reactivity is between chlorine, which is more reactive, and iodine, which is less reactive. It reacts with metals and nonmetals. Phosphorus reacts violently with bromine. Aluminum reacts in a similar way. It can bleach things like chlorine. It reacts better in the light. It dissolves a little in water. Hot water makes it disproportionate into hydrobromic acid and hypobromous acid. Bromide can form compounds with substances such as sodium to form sodium bromide.

Chemical compounds

Bromine forms compounds in many oxidation states: -1, +1, +3, +5, and (sometimes) +7. -1 is the most common. It is found as bromide. Bromides are not reactive. Bromides are colorless. They dissolve in water easily. The other compounds are all strong oxidizing agents. Their chlorine cousins are more common. +1 has the hypobromites, which are unstable. +3 has the bromites. +5 has the bromates, which are more common than all the other ones except the bromides. Bromates are strong oxidizing agents and are sometimes added to flour. +7 has the perbromates. Perbromates are very unstable.

Occurrence

Bromine is found as bromide in the ocean and in brine pools. The Dead Sea has very much bromide in it. Bromine is not an important part of our body, unlike its relatives chlorine and iodine.

Preparation

Bromine is made by bubbling chlorine gas through a solution of a bromide. The bromide is oxidized to bromine, while the chlorine is reduced to chloride.

Uses

It is mainly used to make organobromine compounds, organic compounds with bromine in them. Organobromines are used to put out fires. They used to be added to gasoline. Some were used as pesticides. Some inorganic bromides were used as sedatives. Bromine can also be used as a disinfectant. Silver bromide is used in film.

Safety

Bromine is toxic and corrosive to skin. Bromine gas is irritating when it is breathed in, too. Bromine can react violently with many things.


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